ELECTRONS IN ATOMS
Atomic Structure:
Nucleus:
contains p+ and nº, heavy dense
center of the atom, + charge
NUCLEUS
1
ELECTRONS



Orbit around the nucleus
Negatively charge
Almost no mass
Energy levels (“shells”):
are numbered from the nucleus out (inside out). Each
consecutive energy level has a larger circumference
so it can fit more electrons. (like a stadium)
another name for an energy level is principle quantum number
n = principle quantum number or energy level
# of electrons in the energy level = 2(n2)
energy level:
# of electrons:
1
)2
2
3
4
)8
)18
)32
*outermost energy level can hold a max of 8eStable Octet
–
8 electrons in outermost energy level
-most stable electron configuration because it’s the state
of least energy.
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Metals:
lose e-, have 1,2 or 3e- in the outer shell
Lose 1,2 or 3e- and then the shell underneath has 8e- (stable octet)
*exceptions: (Li, Be and B) –
Non-metals:
lose 1,2 or 3 e- but their shell
underneath is the first energy
level, which can hold a max of 2e-,
and is therefore stable.
gain e-, have 5,6 or 7 e- in their outer shell
gain 3,2 or 1 e- to achieve a stable octet, (8e-)
*exception:
hydrogen – is a gas (not a metal), usually loses 1e- to form H+ ion.
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Subdivisions of Energy Levels in Atoms:
(telephone # with extension) - (609) 268-4600 ext. 2667
Principle quantum # = Energy Levels (1,2,3,4) (period # on table = # of energy levels)
Sublevels (s,p,d,f)
↓
Orbitals
↓
2 electrons each max
↓
Opposed Axial Spins
Pauli Exclusion Principle:
no more than 2e- can occupy an orbital
These electrons must have opposite spins
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Energy levels
are subdivided into sublevels
Energy Level # = # of sublevels
Energy level
Sublevels
1
s
2
s,p
1
3
s,p,d
s)
4
s,p,d,f
Sublevels are divided into
2
sp)
3
sp)d
4
sp)df
orbitals___.
Each sublevel has a different # of orbitals. s has 1; p has 3; d has 5; and f has 7
Orbital-
3 dimensional region of space with a high probability of finding
an electron. (NOT AN ORBIT) Each orbital can hold 2e-max.
Think of a fly trapped in a jar!
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Sublevel
s
*p
d
f
Orbitals
1
Max. # of electrons
2
6
10
14
3
5
7
*p orbitals are designated as
px,py,pz
Hund’s rule: When electrons occupy orbitals of equal energy (entering orbitals of
the same sublevel) they enter one at a time until each orbital of a sublevel contains
1e- each, spinning in the same direction, and then they begin to double up. When
they double up the electrons spin in opposite directions.
↑
↑
px
↑↓
Px
Px
↑
↑
y
z
↑
↑
↑
y
z
↑↓
↑↓
y
z
Px
y
↑↓
↑↓ ↑↓ ↑
Px
↑
y
z
Px
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Electron Configurations of the
Elements #1 - 4
1 H
orbital notation
1s1
electron configuration
2 He
1s2
3 Li
1s2
2s1
1s2
2s2
4 Be
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Physics 2000 Link
TOC → Elements as atoms →
Beyond Hydrogen → Periodic Table
Dave's Whizzy Periodic Table
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Electron Configurations of the Elements #6 - 10
5 B
1s2
2s2
2p x1 y0
z0
1s2
2s2
2p x1 y1
z0
1s2
2s2
2p x1 y1
z1
1s2
2s2
2p x2 y1
z1
1s2
2s2
2p x2 y2
z1
1s2
2s2
2p x2 y2
z2
6 C
7 N
8 O
9 F
10 Ne
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Complete the following:
11 Na
1s2
2s2
2p x2 y2
z2
3s1
1s2
2s2
2p x2 y2
z2
3s2
1s2
2s2
2p x2 y2
z2
3s2
3p x1 y0
z0
1s2
2s2
2p x2 y2
z2
3s2
3p x1 y1
z0
1s2
2s2
2p x2 y2
z2
3s2
3p x1 y1
z1
1s2
2s2
2p x2 y2
z2
3s2
3p x2 y1
z1
1s2
2s2
2p x2 y2
z2
3s2
3p x2 y2
z1
1s2
2s2
x2
z2
3s2
3p x2 y2
z2
12 Mg
13 Al
14 Si
15 P
16 S
17 Cl
18 Ar
2p
y2
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Ar is a noble gas. The 3rd energy level is its outermost and it has a stable octet with
2e- in 3s and 6e- in 3p.
)1
)2
)3
)4
s
sp
spd
spdf
Electrons fill an atom from the nucleus out.
Energy levels are NOT equally spaced. As
you get further from the nucleus, energy
levels get closer together and overlap.
Aufbau Principle: electrons enter orbitals of
Ex.:
lowest energy first
4s fills before 3d, so 4s is of lower energy than the 3d sublevel, even though
the 4s orbital is part of a higher energy level!
Note: After the 3p sublevel is filled, the 4s sublevel begins to fill because this is
where the overlap of energy levels begins!
*Energy levels are NOT evenly spaced.
As you go further out from the
nucleus, the energy levels get closer together and overlap more and more!
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Because electrons “want” to be in the orbital with the LOWEST energy, (remember
the AUFBAU PRINCIPLE), there are some EXCEPTIONS to how sublevels fill up:
Sublevel Energy
Highest
partly full
½ full
full
Lowest
Chromium (Cr)
empty
4s12
Molybdenium (Mo)
3d54
5s12
Copper (Cu)
4d54
4s12
Silver (Ag)
9
3d10
5s12
9
4d10
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Order of Filling Guide
1s
2s
2p
3s
3p
3d
4s
4p
4d
4f
5s
5p
5d
5f
6s
6p
6d
6f*
7s
7p*
7d*
7f*
*Although these sublevels exist in theory, there is no practical use for
them, since all naturally occurring, stable elements are accounted for up
through 6d.
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Bonding: When elements bond (combine to form compounds), only the electrons in
Lewis Dot Diagrams
the outer shell (energy level) are involved. Chemists use _____________________
to show these outer shell electrons.
Lewis Dot Diagrams: Take into account the electrons from the highest energy
levels’ s and p sublevels. Write the symbol for the element and use dots to
represent electrons. Place the dots one on each side before you pair them up
(just like filling orbitals).
Ex.:
H
1s1
H
(1)
Be
C
1s2 2s2
1s2 2s2 2p2
Be
C
(2)
Bonding
electrons
(4)
(2)
S
Ne
1s2 2s2 2p6 3s2 3p4
S
1s2 2s2 2p6
Ne
Paired electrons do
NOT bond
NO bonding electrons
(STABLE OCTET)
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All TRANSITION metals and INNER TRANSITION metals have a
completed s sublevel, so their Lewis Dot Diagrams all look the same:
Ex.:
Fe
1s2 2s2 2p6 3s2 3p6 4s2 3d6
Fe
Pm
1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s2 4d10 5p6 6s2 4f4
Pm
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Use of the Periodic Table to Write Electron Configurations for Elements
d block – group “B” transition
metals; middle of table (10e-)
p block – last 6 columns on
table (6e-)
f block – bottom block on table; “inner transition” metals (14e-)
s block – 1st & 2nd
column on table
(2e-)
(Designate energy level overlaps on your periodic chart)
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Use your periodic chart (or the Order of Filling Chart) to write Electron
Configurations for each of the following:
Abbreviated Form:
Ni # 28 1s2 2s2 2p6 3s2 3p6 4s2 3d8
[Ar] 4s2 3d8
I # 53 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s2 4d10 5p5
[Kr] 5s2 4d10 5p5
Au # 79 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s2 4d10 5p6 6s2 4f14 5d9
[Xe] 6s2 4f14 5d9
2n2 rule: (use “n” to indicate the number of an energy level)
n
n
This rule is used to determine the number of sublevels, number of orbitals,
or number of total electrons in ANY GIVEN ENERGY LEVEL!
How many sublevels, orbitals
=
energy level # [row of periodic table]
& e- could be contained in the
8th energy level?
=
# of sublevels
n2 =
# of orbitals
n = 8 (= # energy levels & # sublevels)
2n2 =
# of electrons
n2 = 64 orbitals
2n2 = 128 e- total
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Atomic orbitals - orbital of an atom, a 3-D region of space with a high probability of
finding an e- (fly in a jar).
spherical
s orbitals are _______________
2-lobed
x, y, and z axes
p orbitals are _______________
and oriented around the ____________________
4-lobed
d orbitals are _______________
with each lobe located ____________________
in one of the quadrants of the x, y, and z axes
__________________________________________________________________
5-lobed
p or d orbitals
f orbitals are _______________
and more complex than ____________________
(f orbitals do not get involved ion bonding)
ORBITAL SHAPES
s orbital
p orbital
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d orbital
Note:
1.
At higher energy levels, the lobes are the same shape, but
farther from the nucleus.
2. When all orbitals of a sublevel are filled, the resulting shape is
spherical.
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Light and the Quantum Mechanical Model
Light
Is made up of Waves
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WAVES
Amplitude Wave’s height from zero to
the crest.
Wavelength
λ
Frequency
ν
Distance between the crests;
crest to crest. Unit meters (m)
Number of wave cycles to pass
A given point per unit of time.
Unit is the Hertz (Hz)
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The wavelength and frequency of light are
Inversely proportional to each other.
As the wavelength of light increases,
the frequency decreases and vice versa.
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Atomic Spectra
Ground State Lowest possible energy of an electron
When electrons absorb energy they can
jump from ground state to an excited state
(a higher energy level).
When the electron drops down to a lower energy
level a quantum (bundle) of energy is released in
the form of light.
Each quantum produces a specific amount of
energy that corresponds to the lines shown on
an atomic spectrum. Each element has its own
unique atomic spectrum.
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